Defining the Atom

1. Defining the Atom • All matter is composed of atoms • Atoms are the smallest particle of an element that retains its identity in a chemical reaction. • Atom comes from the Greek word atomos meaning indivisible. Powerpoint hosted on www.worldofteaching.com Please visit for 100’s more free powerpoints

2. Democritus • Democritus (460 BC – 370 BC) was among the 1st to suggest the existence of atoms. • Democritus believed that atoms were indivisible and indestructible. • His approached was not based on scientific method and was not accepted until later scientific theory.

3. Dalton’s Atomic Theory • The modern atomic thought began with John Dalton (1766-1844) • Dalton used experimental method to transform Democritus’s idea of atoms into a scientific theory. • Dalton studied the ratios in which elements combine and the result was Dalton’s atomic theory.

4. Dalton’s Atomic Theory • The ideas of Dalton’s Theory are: • All elements are composed of tiny particles called atoms. • Atoms of the same element are identical. Atoms of different elements are different. • Atoms of different elements can physically mix together or chemically combine in simple whole number ratios to form compounds. • Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element are never changed.

5. Sizing Up Atoms • ATOMS ARE TINY!!!!!! • A pure copper coin the size of a penny contains 2.4 x 1022 copper atoms. There are only about 6 x 109 people on Earth!!! • Individual atoms are observable with instruments such as scanning tunneling microscopes.

6. Subatomic Particles • Atoms are now know to be broken down into smaller, more fundamental particles called subatomic particles. • There are 3 kinds of subatomic Particles. • Electrons • Protons • Neutrons

7. Electrons • Discovered in 1897 by J.J. Thomson (1856-1940) an English physicist. • They are negatively charged subatomic particles. • Thomson performed experiments using a sealed glass tube with gases in it. He passed an electric current through the tube and the result was a cathode ray.

8. Electrons Continued • U.S. physicist Robert A. Millikan (1868-1953) carried out experiments to find the quantity of charge carried out by an electron. • He calculated the mass of the electron; that is very similar to the excepted mass given today.

9. Protons • Eugen Goldstein (1850-1930) observed in 1886 that in a cathode-ray tube there were rays going in the opposite direction. He concluded they were positively charged particles. • Protons are positively charged subatomic particles. • Each proton has a mass about 1840 times that of an electron.

10. Neutrons • In 1932 English physicist James Chadwick (1891-1974) confirmed the existence of another subatomic particle. • Neutrons are subatomic particles with no charge but a mass nearly equal to the proton.

11. Rutherford’s Gold-Foil Experiment • This experiment change the prevailing thought of the structure of the atom. • Because of this experiment Rutherford concluded that most the of alpha particles pass through b/c the atom is mostly empty space, and the reason that some of the particles were deflected were due to the concentration of the positive charge in the atom which is now known as the nucleus.

12. Structure of the atom • The nucleus of the atom is the tiny central core of an atom and is composed of protons and neutrons. • Rutherford’s model is the nuclear atom: In the nuclear atom, the proton and neutron are located in the nucleus. The electrons are distributed around the nucleus. • This will later be revised, let’s sneak peak @ page 128.

13. Distinguishing Among Atoms • We learned yesterday that atoms are composed of protons, neutrons, and electrons. • Protons and Neutrons are in the nucleus and electrons surround the nucleus. • So how do elements differ if they all contain atoms?

14. Atomic Number • Elements are different because they contain different numbers of protons. • The atomic number of an element is the number of protons in the nucleus of an atom in that element. • Ex: The atomic number of Hydrogen (H) is 1, so all hydrogen atoms have 1 proton. • Ex: All Oxygen atoms (O) have 8 protons, so the atomic number of Oxygen is 8.

15. Atomic Number • Remember all atoms are electrically neutral. • Therefore; the number of Protons equal the number of Electrons. • Meaning the number of negatively charged particles must equal the number of postively charged particles.

16. Mass Number • Most of the mass of an atom is concentrated in its nucleus. • The mass number of an atom is the total number of protons and neutrons in an atom. • Ex: Helium (He) has 2 protons and 2 neutrons: its mass number is 4. • Ex: Carbon (C) has 6 protons and 6 neutrons: its mass number is12.

17. Mass Number • If you know the atomic number and mass number of an atom, you can determine its composition. • The number of neutrons in an atom is the difference between the mass number and atomic number. • # of neutrons = Mass # - Atomic #

18. Mass Number • The composition of an atom can be written in Shorthand notation using atomic number and mass number. • Ex: Look at Figure 4.8 • The atomic number is the subscript and the mass number is the superscript. • How many electrons does a gold atom have? • How many neutrons does a gold atom have?